1. Details of Module and its structure

Module Detail

Subject Name Chemistry

Course Name Chemistry 03 (Class XII, Semester 01)

Module Name/Title Coordination Chemistry: Part 3

Module Id lech_10903

Pre-requisites Stereoisomerism, geometrical and optical isomerism,paramagmetic and diagmagnetic compounds, valence bond andcrystal field theory

Objectives After going through this module, the learner will be able to: Draw relation between the colour of coordination

compounds and the wavelengths of the light in the visiblerange

Understand the special type of bond between metalcarbonyls

Learn the stability of coordination compounds Explain the applications of coordination compounds in

extraction of metals, medicine and qualitative analysis

Keywords Coloured Complexes, Visible Range, ElectromagneticRadiation, Synergic Bonding, Stability Constant

2. Development Team

Role Name Affiliation

National MOOC Coordinator (NMC)

Prof. Amarendra P. Behera CIET, NCERT, New Delhi

Program Coordinator Dr. Mohd. Mamur Ali CIET, NCERT, New Delhi

Course Coordinator (CC) / PI Prof. Alka MehrotraProf. R. K. Parashar

DESM, NCERT, New DelhiDESM, NCERT, New Delhi

Course Co-Coordinator / Co-PI Dr. Aerum Khan CIET, NCERT, New Delhi

Subject Matter Expert (SME) Ms Sarojini Sinha SAJS, Vasundhara, Gzb, UP

Review Team Dr. Sushmita ChoudharyDr. Aerum Khan

Gargi College, University of DelhiCIET, NCERT

Language Editor Dr. Sushmita ChoudharyDr. Aerum Khan

Gargi College, University of DelhiCIET, NCERT

Table of Contents:

1. Introduction

2. Colour in Coordination Compounds

3. Bonding in Metal Carbonyls

4. Stability of Coordination Compounds

5. Importance and Applications of Coordination Compounds/

6. Summary

1. Introduction

Visible light is a very small portion of the electromagnetic spectrum that human eyes are sensitive

to. It has a range of wavelengths that our eyes can detect. The colour of most objects depends upon

the interaction between visible light and the electrons of atoms or molecules that make up the

object. It is usually the result of a dynamic process on the molecular level: the absorption of light

and the resulting change of a molecule's quantized energy. This absorption coloration mechanism is

responsible for the colours of grass, blood etc. but not of the sky. The sky's colour is due to selective

scattering of different wavelengths of sunlight by the molecules of nitrogen and oxygen in the

atmosphere.

The perceived colour of an object has a complementary relationship with the colour of the visible

light absorbed. First consider white light, a mixture of all visible wavelengths, impinging on a

coloured object. The object absorbs some wavelengths of the light, the wavelengths of light that are

not absorbed are transmitted or reflected to the observer's eye. A substance that appears blue is

transmitting or reflecting blue light to the eye and absorbing other colors of the white light that are

not blue, i.e. It would absorb red, orange, yellow, and violet light. The absorption spectrum would

show high absorbance of all visible wavelengths, besides blue. The transmission or reflectance

spectrum would have a maximum at a wavelength corresponding to blue light.

2. Colour in Coordination Compounds

One of the most distinctive properties of transition metal complexes is their wide range of colours.

This means that some of the visible spectrum is being removed from white light as it passes through

the sample, so the light that emerges is no longer white. The colour of the complex is

complementary to that which is absorbed. The complementary colour is the colour generated from

the wavelength left over; if green light is absorbed by the complex, it appears red. Table 1 gives the

relationship of the different wavelength absorbed and the colour observed.

The colour in the coordination compounds can be readily explained in terms of crystal field theory.

Consider, for example, the complex [Ti(H2O)6]3+, which is violet in colour. This is an octahedral

complex where the single electron (Ti3+ is a 3d1 system) in the metal d orbital is in the t2g level in the

ground state of the complex. The next higher state available for the electron is the vacant eg level. If

light corresponding to the energy of blue-green region is absorbed by the complex, it would excite

the electron from t2g level to the eg level (t2g1 eg

0 → t2g0 eg

1). Consequently, the complex appears

violet in colour (Fig. 1). The crystal field theory attributes

the colour of the coordination compounds to d-d transition of the electron.

Fig 1: Transition of an electron in [Ti(H2O)6]3+

It is important to note that in the absence of ligand, crystal field splitting does not occur and hence

the substance is colourless. For example, removal of water from [Ti(H2O)6]Cl3 on heating renders it

colourless. Similarly, anhydrous CuSO4 is white, but CuSO4.5H2O is blue in colour. The influence

of the ligand on the colour of a complex may be illustrated by considering the [Ni(H2O)6]2+

complex, which forms when nickel(II) chloride is dissolved in water. If the didentate ligand, ethane-

1,2-diamine(en) is progressively added in the molar ratios en:Ni, 1:1, 2:1, 3:1, the following series

of reactions and their associated colour changes occur:

[Ni(H2O)6]2+ (aq) + en (aq) = [Ni(H2O)4(en)]2+ (aq) + 2H2O

green pale blue

[Ni(H2O)4 (en)]2+(aq) + en (aq) = [Ni(H2O)2(en)2]2+ (aq) + 2H2O

blue/purple

[Ni(H2O)2(en)2]2+ (aq) + en (aq) = [Ni(en)3]2+ (aq) + 2H2O

violet

This sequence is shown in Fig. 2.

Fig 2: Aqueous solutions of complexes of nickel(II) with an increasing number of ethane-1,

2-diamine ligands

Colour of Some Gem Stones

The colours produced by electronic transitions within the d orbitals of a transition metal ion

occur frequently in everyday life. Ruby [Fig.3 a] is aluminium oxide (Al2O3) containing about

0.5-1% Cr3+ ions (d3), which are randomly distributed in positions normally occupied by Al3+ .

We may view these chromium(III) species as octahedral chromium(III) complexes

incorporated into the alumina lattice; d–d transitions at these centres give rise to the colourIn

emerald [Fig.3 b], Cr3+ ions occupy octahedral sites in the mineral beryl (Be3Al2Si6O18). The

absorption bands seen in the ruby shift to longer wavelength, namely yellow-red and blue,

causing emerald to transmit light in the green region.

Limitations of Crystal Field Theory

Intext Questions:

Problem 1: Explain on the basis of valence bond theory that [Ni(CN)4]2– ion with square planar

structure is diamagnetic and the [NiCl4]2– ion with tetrahedral geometry is paramagnetic.

Solution:Nickel is in +2 oxidation state in[Ni(CN)4]2– complex. The formation of [Ni(CN)4]2– and

[NiCl4]2–can be explained on basis of hybridization as below:

dsp2 hybrid orbitals accommodate four pairs of electrons from four CN— groups and the resulting

square planar complex is diamagnetic since it does not contain any unpaired electrons. Ideally the

complex should be following sp3hybridisation but it is actually dsp2, ie; the 2 unpaired 3d electrons

pair up against Hund’s rule before hybridization. This is on account of the following reasons:

a) If no pairing occurs the complex should be paramagnetic but it was actually found to be

diamagnetic.

b) CN—is a strong ligand and as it approaches the metal ion, the electrons pair up.

On the other hand in [NiCl4]2–Cl—provides a weak ligand field. It is therefore, unable to cause a

pairing of the 3d unpaired electrons in Ni2+ ion. Hence, sp3hybridisation involving 4s and 4p

orbitals takes place and the shape of the molecule is tetrahedral. This is paramagnetic in nature.

Problem 2: [NiCl4]2– is paramagnetic while [Ni(CO)4] is diamagnetic though both are tetrahedral.

Why?

Solution: In [NiCl4]2–, Ni is in +2 oxidation state with the configuration 3d8 4s0 and Cl—provides a

weak ligand field. It is therefore, unable to cause a pairing of the 3d unpaired electrons in Ni2+ ion.

Hence, sp3hybridisation involving 4s and 4p orbitals takes place and the shape of the molecule is

tetrahedral. This is paramagnetic in nature. While in [Ni(CO)4], Ni is in zero oxidation state with the

configuration 3d8 4s2. In the presence of CO ligand which is a strong field ligand, the 4s electrons

shift to 3d orbitals and pair up, thus there is no unpaired electron present in the molecule hence it is

diamagnetic in nature. The formation of [Ni(CO)4]complex can be explianed as below:

Problem 3:[Fe(H2O)6]3+ is strongly paramagnetic whereas [Fe(CN)6]3– is weakly paramagnetic.

Explain.

Solution: In both the complexes Fe is in +3 oxidation state with the configuration 3d5, CN—is a

strong field ligand, hence in its presence the 3d unpaired electrons pair up leaving only 1 unpaired

electron. The hybridization is d2sp3 forming an inner orbital complex. H2O is a weak field ligand

and hence in its presence the 3 d electrons do not pair up and the hybridization is sp3d2 forming an

outer orbital complex with 5 unpaired electrons. Hence it is strongly paramagnetic. Formation of

[Fe(H2O)6]3+ and [Fe(CN)6]3– is shown below:

Problem 4:Explain [Co(NH3)6]3+ is an inner orbital complex whereas [Ni(NH3)6]2+ is an outer

orbital complex.

Solution: In [Co(NH3)6]3+, Co is in +3 oxidation state with the configuration 3d6. In presence of NH3

3d electrons pair up leaving two orbitals of 3d empty. Hence the hybridisation is d2sp3 forming an

inner orbital complex. In [Ni(NH3)6]2+, Ni is in +2 oxidation state with the configuration 3d8. In

presence of NH3, the 3d electrons do not pair up. The hybridization involved is sp3d2 forming an

outer orbital complex.

Problem 5: Predict the number of unpaired electrons in the square planar [Pt(CN)4]2– ion.

Solution: 78Pt lies in group 10 with the configuration 5d9 6s1. Hence, Pt2+ has the configuration 5d8.

For square planar shape, the hybridization is dsp2. Hence, the unpaired electrons in 5d pair up and

thus there is one d-orbital empty, hence the hybridization dsp2 is possible. Therefore there are no

unpaired electrons in [Pt(CN)4]2– ion.

Problem 6: The hexaquomanganese(II) ion contains five unpaired electrons, while the

hexacyanoion contains only one unpaired electron. Explain using Crystal Field Theory.

Solution: The electronic configuration of Mn in +2 oxidation state is, 3d5. In presence of H2O

ligand, the distribution of these five electrons becomes as; 3 electrons in 3-t2g orbitals and 2

electrons in 2-eg orbitals (t2g3and eg

2). Thus, all electrons remain unpaired. In the presence of CN—

ligand, the distribution becomes as 5 electrons in 3-t2g orbitals, therefore the pairing of electrons

takes place with 5 electrons occupying 3-t2g orbitals and none of the electrons occupy the 2-eg

orbitals, therefore 2-eg orbitals remain empty. One t2g orbital now has an unpaired electron while the

other 2 t2g orbitals are completely filled with 2 electrons in each orbital.

3. Bonding in Metal Carbonyls

Metal carbonyls are a class of compounds involving carbon monoxide (CO) as a ligand. In

homoleptic carbonyls the metal is attached to only CO ligands, the metal is generally inzero

oxidation state. Tetracarbonylnickel(0), pentacarbonyliron (0) and hexacarbonylchromium(0) have

tetrahedral, trigonalbipyramidal and octahedral structures respectively.Polynuclear carbonyls are

also known. Decacarbonyldimanganese (0), Mn2(CO)10 has two Mn(CO)5 units linked by a Mn –

Mn bond. In octacarbonyldicobalt(0), Co2(CO)8 there is a Co – Co bond and two carbonyl groups

act as bridge between the two Co atoms (Fig. 4).

The metal_carbon bond in metal carbonyls possesses both sigma and pi bond character.

The M-C sigma bond is formed by donation of a lone pair of electrons on the carbonyl carbon to the

vacant orbitals of the metal. The M-C pi bond is formed by donation of a pair of electrons from the

filled d orbital of the metal into the vacant antibonding orbitals of the carbon monoxide. The metal

ligand bonding creates a synergic effect which strengthens the bond between CO and the metal (Fig.

5).

Fig. 4 Structures of some representative homoleptic metal carbonyls

Fig. 5: Synergic Bonding

4. Stability of Coordination Compounds

The stability of a complex in solution refers to the degree of association between the two species

involved in the state of equilibrium. The magnitude of the (stability or formation) equilibrium

constant for the association, quantitatively expresses the stability. Thus, if we have a reaction of the

type:

M + 4L -------- ML4

then the greater the stability constant, the higher is the proportion of ML4 that exists in solution.

Free metal ions rarely exist in the solution so that M will usually be surrounded by solvent

molecules which will compete with the ligand molecules, L, and be successively replaced by them.

For simplicity, we generally ignore these solvent molecules and write four stability constants as

follows:

M + L --------- ML K1 = [ML]/[M][L] ML + L ------------- ML2 K2 = [ML2 ]/[ML][L] ML2 + L --------- ML3 K3 = [ML3 ]/[ML2 ][L] ML3 + L ----------- ML4 K4 = [ML4 ]/[ML3 ][L]

where K1 , K2 , etc., are referred to as stepwise stability constants. Alternatively, we can write the

overall stability constant thus:

M + 4L -------- ML4 β4 = [ML4]/[M][L]4

The stepwise and overall stability constant are therefore related as follows:

β4 = K1 × K2 × K3 × K4 or more generally,

βn = K1 × K2 × K3 × K4 ....... Kn

If we take as an example, the steps involved in the formation of the tetraamminecopper (II) ion, we

have the following:

Cu2+ + NH3----------Cu(NH3)2+ K1=[Cu(NH3)]2+/[Cu2+][NH3] Cu(NH3)2+ + NH3--------[Cu(NH3)2]2+ K2=[Cu(NH3)2]2+/[Cu(NH3)]2+[NH3]

where K1, K2 are the stepwise stability constants. The overall stability constant β4 is represented as

β4 = [Cu(NH3)4]2+/[Cu2+][NH3]4

The addition of the four ammonia molecules to copper shows a pattern found for most formation

constants, in that the successive stability constants decrease. In this case, the four constants are:

logK1 = 4.0, logK2 = 3.2, logK3 = 2.7, logK4 = 2.0 or log β4 = 11.9

The instability constant or the dissociation constant of coordination compounds is defined as the

reciprocal of the formation constant.

Intext Question

Problem 7:Calculate the overall complex dissociation equilibrium constant for the [Cu(NH3)4]2+

ion, given that β4 for this complex is 2.1 × 1013 .

Solution: The Overall stability constant (β4) = 2.1 X 1013

The overall dissociation constant is the reciprocal of the overall stability constant. Hence, overall

complex dissociation constant = 1

β 4=

1

2.1 X 1013=4.7 X 10− 14

5. Importance and Applications of Coordination Compounds

The coordination compounds are of great importance. These compounds are widely present in the

mineral, plant and animal worlds and are known to play many important functions in the area of

analytical chemistry, metallurgy, biological systems, industry and medicine.

These are described below:

Coordination compounds find use in many qualitative and quantitative chemical analyses.

The familiar colour reactions given by metal ions with a number of ligands (especially

chelating ligands), as a result of formation of coordination entities, form the basis for their

detection and estimation by classical and instrumental methods of analysis. Examples of

such reagents include EDTA, DMG (dimethylglyoxime), α–nitroso–β–naphthol, cupron, etc.

Hardness of water is estimated by simple titration with Na2EDTA. The Ca2+ and Mg2+ ions

form stable complexes with EDTA. The selective estimation of these ions can be done due

to difference in the stability constants of calcium and magnesium complexes.

Some important extraction processes of metals, like those of silver and gold, make use of

complex formation. Gold, for example, combines with cyanide in the presence of oxygen

and water to form the coordination entity [Au(CN)2]– in aqueous solution. Gold can be

separated in metallic form from this solution by the addition of zinc.

Similarly, purification of metals can be achieved through formation and subsequent

decomposition of their coordination compounds. For example, impure nickel is converted to

[Ni(CO)4 ], which is decomposed to yield pure nickel.

Coordination compounds are of great importance in biological systems. The pigment

responsible for photosynthesis, chlorophyll, is a coordination compound of magnesium.

Haemoglobin, the red pigment of blood which acts as oxygen carrier is a coordination

compound of iron. Vitamin B12, cyanocobalamine, the anti– pernicious anaemia factor, is a

coordination compound of cobalt. Among the other compounds of biological importance

with coordinated metal ions are the enzymes like, carboxypeptidase A and carbonic

anhydrase (catalysts of biological systems).

Coordination compounds are used as catalysts for many industrial processes. Examples

include rhodium complex, [(Ph3P)3RhCl], a Wilkinson catalyst, is used for the

hydrogenation of alkenes.

Articles can be electroplated with silver and gold much more smoothly and evenly from

solutions of the complexes, [Ag(CN)2]– and [Au(CN)2]– than from a solution of simple metal

ions.

In black and white photography, the developed film is fixed by washing with hypo solution

which dissolves the undecomposedAgBr to form a complex ion, [Ag(S2O3)2]3.

There is growing interest in the use of chelate therapy in medicinal chemistry. An example is

the treatment of problems caused by the presence of metals in toxic proportions in

plant/animal systems. Thus, excess of copper and iron are removed by the chelating ligands

D–penicillamine and desferrioxime B via the formation of coordination compounds. EDTA

is used in the treatment of lead poisoning. Some coordination compounds of platinum

effectively inhibit the growth of tumours. Examples are: cis–platin and related compounds.

6. Summary

The chemistry of coordination compounds is an important and challenging area of modern

inorganic chemistry. During the last fifty years, advances in this area, have provided development of

new concepts and models of bonding and molecular structure, novel breakthroughs in chemical

industry and vital insights into the functioning of critical components of biological systems. The

first systematic attempt at explaining the formation, reactions, structure and bonding of a

coordination compound was made by A. Werner. His theory postulated the use of two types of

linkages (primary and secondary) by a metal atom/ion in a coordination compound. In the modern

language of chemistry these linkages are recognised as the ionisable (ionic) and non-ionisable

(covalent) bonds, respectively. Using the property of isomerism, Werner predicted the geometrical

shapes of a large number of coordination entities. The Valence Bond Theory (VBT) explains with

reasonable success, the formation, magnetic behaviour and geometrical shapes of coordination

compounds. It, however, fails to provide a quantitative interpretation of magnetic behaviour and has

nothing to say about the optical properties of these compounds. The Crystal Field Theory (CFT) to

coordination compounds is based on the effect of different crystal fields (provided by the ligands

taken as point charges), on the degeneracy of d orbital energies of the central metal atom/ion. The

splitting of the d orbitals provides different electronic arrangements in strong and weak crystal

fields. The treatment provides for quantitative estimations of orbital separation energies, magnetic

moments and spectral and stability parameters. However, the assumption that ligands consititute

point charges creates many theoretical difficulties. The metal–carbon bond in metal carbonyls

possesses both σ and π character. The ligand to metal is σ bond and metal to ligand is π bond. This

unique synergic bonding provides stability to metal carbonyls. The stability of coordination

compounds is measured in terms of stepwise stability (or formation) constant (K) or overall stability

constant (β). The stabilisation of coordination compound due to chelation is called the chelate

effect. The stability of coordination compounds is related to Gibbs energy, enthalpy and entropy

terms. Coordination compounds are of great importance. These compounds provide critical insights

into the functioning and structures of vital components of biological systems. Coordination

compounds also find extensive applications in metallurgical processes, analytical and medicinal

chemistry.